AQA GCSE Chemistry 8462: Chemical changes – Revision Notes

These are condensed recall notes for section 4.4 Chemical changes (4.4.1.1 to 4.4.3.5) of the AQA GCSE Chemistry (8462) specification, for teaching from September 2016 and exams from June 2018 onwards. The topic is assessed on Paper 1, set at Foundation and Higher Tier. Points the specification marks (HT only) are labelled Higher tier only. Required practicals 1, 2 and 3 are in this topic.

For full explanations and worked examples, use the Chemical changes study guide. Then test yourself with the Chemical changes practice questions. The course hub is AQA GCSE Chemistry, and the printable checklist lists every point. Titration maths relies on Quantitative chemistry.

Key definitions

Term	Meaning	Tier
Oxidation	Gain of oxygen	Both
Reduction	Loss of oxygen	Both
Oxidation	Loss of electrons	Higher tier only
Reduction	Gain of electrons	Higher tier only
Electrolyte	Molten or dissolved ionic compound that conducts; ions free to move	Both
Cathode	Negative electrode; positive ions go here	Both
Anode	Positive electrode; negative ions go here	Both
Strong acid	Completely ionised in aqueous solution	Higher tier only
Weak acid	Only partially ionised in aqueous solution	Higher tier only

4.4.1 Reactivity of metals

Reactivity series (most to least reactive):

potassium  sodium  lithium  calcium  magnesium  [carbon]  zinc  iron  [hydrogen]  copper

• Reactivity depends on how easily a metal forms its positive ion.
• Metal + water → metal hydroxide + hydrogen. Metal + acid → salt + hydrogen. Room temperature only; no steam.
• K, Na, Li fizz in water (K and Na melt; K gives a lilac flame). Ca fizzes. Mg barely reacts with water but fizzes well in acid. Zn and Fe react with acid (Fe slowly) but not visibly with water. Cu reacts with neither.
• A more reactive metal displaces a less reactive metal from its compound.
• Deducing an order: more bubbles or faster temperature rise → more reactive; a metal that displaces another sits above it.

Extraction:

• Gold and other unreactive metals occur as the metal itself.
• Below carbon → reduce the oxide with carbon. The oxide loses oxygen (reduced); carbon gains oxygen (oxidised).
• Above carbon (or reacts with carbon) → electrolysis of the molten compound.

Method in steps: ionic equation for displacement (Higher tier only)

1. Write the full equation: Fe + CuSO₄ → FeSO₄ + Cu.
2. Split the soluble ionic compounds into ions.
3. Cross out the ions unchanged on both sides (spectator ions, here SO₄²⁻).
4. Write what is left: Fe + Cu²⁺ → Fe²⁺ + Cu.
5. Iron loses electrons, so it is oxidised. Copper ions gain electrons, so they are reduced.

4.4.2 Reactions of acids

Salt-making reactions

Reaction	Products
Acid + metal (Mg, Zn, Fe)	Salt + hydrogen
Acid + alkali (soluble hydroxide)	Salt + water
Acid + base (insoluble hydroxide or metal oxide)	Salt + water
Acid + metal carbonate	Salt + water + carbon dioxide

Naming: hydrochloric → chloride, nitric → nitrate, sulfuric → sulfate. The first part comes from the metal.

Ions for formulae:

Positive	Negative
Na⁺, K⁺	Cl⁻, NO₃⁻, OH⁻
Mg²⁺, Ca²⁺, Zn²⁺, Cu²⁺	SO₄²⁻, CO₃²⁻, O²⁻
Al³⁺, Fe³⁺

Balance the charges to zero. Use brackets when more than one polyatomic ion is needed: Ca(NO₃)₂, Al₂(SO₄)₃.

Higher tier only: metal + acid is a redox reaction. Mg + 2H⁺ → Mg²⁺ + H₂: Mg is oxidised, H⁺ is reduced.

Required practical 1 – method in steps

1. Warm dilute acid with a Bunsen burner.
2. Add the insoluble oxide or carbonate in small amounts, stirring, until some is left unreacted (excess).
3. Filter off the excess solid.
4. Evaporate some water from the filtrate using a water bath or electric heater.
5. Leave to cool and crystallise; filter off the crystals and pat dry.

Why excess solid? So all the acid reacts and the salt is not contaminated with acid.

pH and neutralisation

• Acids → H⁺ ions. Alkalis → OH⁻ ions.
• pH 0–14; 7 neutral; below 7 acid; above 7 alkali.
• Universal indicator (or wide range indicator) gives an approximate pH by colour; a pH probe gives a reading.
• Neutralisation: H⁺ + OH⁻ → H₂O.

Required practical 2: titration – method in steps

1. Pipette 25.0 cm³ of alkali into a conical flask; add a few drops of a suitable indicator.
2. Fill the burette with acid; read the start volume at eye level.
3. Add acid, swirling; go drop by drop near the colour change.
4. Stop at the colour change; record the final volume; titre = final − start.
5. Repeat until results agree closely; average those results.

Only strong acids (sulfuric, hydrochloric, nitric) and strong alkalis are used.

Titration calculation (Higher tier only)

25.0 cm³ of KOH is neutralised by 22.00 cm³ of 0.100 mol/dm³ HCl.
HCl + KOH → KCl + H2O  (ratio 1 : 1)
moles HCl = 0.100 × 0.02200 = 0.00220 mol  =  moles KOH
[KOH] = 0.00220 ÷ 0.0250 = 0.0880 mol/dm³
in g/dm³: 0.0880 × 56 = 4.93 g/dm³

Strong and weak acids (Higher tier only)

• Strong: hydrochloric, nitric, sulfuric. Weak: ethanoic, citric, carbonic.
• Same concentration → stronger acid has lower pH.
• Each fall of 1 pH unit → [H⁺] × 10. Two units → × 100. Three units → × 1000.

4.4.3 Electrolysis

Molten compounds

Metal at the cathode, non-metal at the anode (inert electrodes). Molten lead bromide → lead + bromine.

Aluminium

• Electrolyte: aluminium oxide dissolved in molten cryolite → lower melting temperature, so less energy.
• Carbon anode reacts with the oxygen made there → carbon dioxide → anode wears away and is replaced.
• Energy is needed both to melt the compounds and to produce the current.

Aqueous solutions – method in steps

1. Cathode: is the metal more reactive than hydrogen? Yes → hydrogen. No → the metal.
2. Anode: is there a halide ion (Cl⁻, Br⁻, I⁻)? Yes → the halogen. No → oxygen.

Solution	Cathode	Anode
Copper chloride	Copper	Chlorine
Sodium sulfate	Hydrogen	Oxygen
Silver nitrate	Silver	Oxygen
Potassium iodide	Hydrogen	Iodine

The H⁺ and OH⁻ come from water molecules breaking down. Required practical 3 tests these predictions: write a hypothesis first, use inert electrodes, and record what forms at each electrode (bubbles of gas, or a coating of metal on the cathode). Change only the solution each time.

Half equations (Higher tier only)

Electrode	Process	Examples
Cathode (−)	Gain of electrons = reduction	2H⁺ + 2e⁻ → H₂; Cu²⁺ + 2e⁻ → Cu; Al³⁺ + 3e⁻ → Al
Anode (+)	Loss of electrons = oxidation	2Cl⁻ → Cl₂ + 2e⁻; 4OH⁻ → O₂ + 2H₂O + 4e⁻

Check that atoms and charges balance on each side.

Must-know distinctions

• Strong vs concentrated: strength is about ionisation; concentration is about amount of acid per volume.
• Base vs alkali: an alkali is a soluble base.
• Oxygen definition vs electron definition: both tiers use oxygen; only Higher uses electrons.
• Molten vs aqueous: molten gives the metal; aqueous may give hydrogen instead.
• Displacement vs electrolysis: displacement is a reactive metal pushing out a less reactive one; electrolysis uses a current.

Quick self-test

1. Define reduction in terms of oxygen.
2. Put in order, most reactive first: iron, potassium, magnesium, copper.
3. Name the salt made from zinc oxide and nitric acid.
4. Give the formula of calcium nitrate.
5. What are the products of magnesium carbonate + hydrochloric acid?
6. Why is excess solid used when making a soluble salt?
7. Write the ionic equation for neutralisation.
8. (Higher tier only) A solution changes from pH 5 to pH 3. By what factor does [H⁺] change?
9. (Higher tier only) How many moles of HCl are in 20.0 cm³ of 0.50 mol/dm³ acid?
10. Name the products at each electrode when sodium chloride solution is electrolysed.
11. Why is cryolite used in making aluminium?
12. (Higher tier only) Write the half equation for chloride ions at the anode.

Answers

1. Loss of oxygen.
2. Potassium, magnesium, iron, copper.
3. Zinc nitrate (and water).
4. Ca(NO₃)₂.
5. Magnesium chloride, water and carbon dioxide.
6. So that all the acid reacts; the unreacted solid is filtered off.
7. H⁺ + OH⁻ → H₂O.
8. × 100 (increases 100 times).
9. 0.50 × 0.0200 = 0.010 mol.
10. Hydrogen at the cathode; chlorine at the anode.
11. It lowers the melting temperature, so less energy is needed.
12. 2Cl⁻ → Cl₂ + 2e⁻.

Where marks are usually lost

• Saying copper reacts slowly with dilute acid; it does not react.
• Naming carbon as “reduced” in a carbon extraction.
• Writing Ca(NO₃)₂ as CaNO₃ or CaNO₃₂.
• Forgetting carbon dioxide in the products of an acid + carbonate reaction.
• In required practical 1, heating the solution until dry instead of leaving it to crystallise.
• In a titration description, leaving out swirling, drop-by-drop addition near the end point or repeating.
• (Higher tier only) Describing a weak acid as “dilute”.
• (Higher tier only) Saying one pH unit changes [H⁺] by 1 rather than by a factor of 10.
• Predicting sodium at the cathode for sodium chloride solution.
• (Higher tier only) Writing electrons on the wrong side of a half equation, so charges do not balance.

Official syllabus

AQA GCSE Chemistry (8462) specification, Version 1.1 (October 2019), for teaching from September 2016 and exams from June 2018 onwards (AQA), section 4.4 Chemical changes. Check your recall with the free 10-minute diagnostics.